The present invention relates to the field of measurements and in particular to apparatus and methods for optically measuring the state of charge of at least one electrolyte.
The present invention relates generally to reduction-oxidation, or redox flow batteries used to store electrical energy in a chemical form, and subsequently dispense the stored energy in an electrical form via a spontaneous reverse redox reaction. Aspects of redox flow batteries incorporating external liquid electrolytes have been described (see for example Thaller, U.S. Pat. No. 3,996,064, herein incorporated by reference).
A redox flow battery is an electrochemical storage device in which an electrolyte containing one or more dissolved electroactive species flows through a reactor converting chemical energy to electrical energy and vice versa. Additional electrolyte is stored externally, (for example in tanks) and flows through a set of cells (e.g. by pumps, gravity or other movement) where the electrochemical reaction takes place. Typically, the reaction in a flow battery is reversible, i.e. it can be recharged without replacing the electroactive material. The energy capacity of a redox flow battery is related to the electrolyte volume (tank size). The discharge time of a redox flow battery at full power varies from several minutes to many days. Advantages of flow cells over standard fuel cells and batteries may include some or all of flexible device layout (due to separation of the power and energy components), long cycle life (because there are no solid-solid phase changes), no harmful emissions are generated, low maintenance and tolerance to overcharge/overdischarge. Disadvantages may include their complicated apparatus, (requiring pumps, sensors, control units, secondary containment vessels, etc) and low energy densities.
A flow battery differs from a secondary battery in the sense that flow batteries maintain the majority of the electrolyte outside of the cell volume, pumping it into the cell as needed. Hence the power and energy capacity are decoupled. Secondary batteries retain all of the electrolyte within the cell volume, and hence the power and energy capacity are coupled. However, both a flow battery and a secondary battery are rechargeable.
A flow battery differs from a fuel cell in the sense that although both work on electrochemical redox principles, in the latter a fuel is generally consumed and the system is generally not rechargeable. Conventional fuel cell fuels include hydrogen, methanol, gasoline, etc. The fuel must be continuously replenished in order to produce power. The electrolytes in a flow battery are rechargeable, and therefore an external fuel supply is unnecessary.
It may be noted that the minimal unit that performs electrochemical energy conversion is generally called a “cell”, whether in the case of flow batteries, fuel cells or secondary batteries. A device that integrates many such cells, in series or parallel, to obtain higher current or voltage or both, is generally referred to as a “battery”. However, it is common to refer to a single cell used on its own as a battery rather than a cell, thus leading to some possible confusion.
The redox flow cell works by changing the oxidation state of its constituents during charging or discharging. The basic cell consists of two half-cells, connected in series by the conductive electrolyte, one for anodic reaction and the other for cathodic reaction. Each half-cell comprises an electrode with a defined surface area upon which the redox reaction takes place. Electrolyte flows through the half-cell as the redox reaction takes place. The two half-cells are separated by an ion-exchange membrane that allows primarily either positive ions or negative ions to pass through it. Multiple such cells can be stacked either in series to achieve higher voltage or in parallel in order to achieve higher current. The reactants are stored in separate tanks and dispensed into the cells as necessary in a controlled manner.
A non limiting, illustrative example of a Redox pair would include:Fe3++e−→Fe2+(Eo=+0.771V)Cr3++e−→Cr2+(Eo=−0.407V)
where Eo is the standard electrode potential of the reaction.
If the electrolyte has a net higher positive electrode potential (Eo) compared to a Standard Hydrogen Electrode (SHE) during discharge of the system, then the electrolyte is called the catholyte. The complementary electrolyte is then called the anolyte.
In a simple implementation of the redox cell technology, an acidic solution of FeCl2 is on the cathode side and an acidic solution of CrCl3 is on the anode side. Upon applying an appropriate positive voltage on the cathode with respect to the anode, the following reactions take place:Cathodic reaction: Fe2+→Fe3++e−Anodic reaction: Cr3++e−→Cr2+
Applying the external power supply affects an electron transfer, while a Cl− ion crosses the membrane from the anodic half-cell to the cathodic half-cell through the ion exchange membrane in order to preserve the charge balance. In the ideal situation, the fully charged flow cell consists of 100% FeCl3 solution on the cathode side and 100% CrCl2 solution on the anode side.
When the external power supply is replaced with a load, the cell begins to discharge, and the opposite Redox reactions take place:Cathodic reaction: Fe3++e−→Fe2+Anodic reaction: Cr2+→Cr3++e−
Therefore, in the most ideal situation, the fully discharged flow cell consists of 100% FeCl2 solution on the cathode side and 100% CrCl3 solution on the anode side.
A variation of the Fe/Cr system described above is a redox cell with premixed Fe and Cr solutions (see Gahn et al, NASA TM-87034). Since no membrane is perfectly perm-selective, anolyte and catholyte eventually become cross-mixed over many cycles of charge and discharge, thus reducing the net system capacity. Gahn et al proposed a remedy to this problem using a redox cell, both sides of which contain FeCl2 and CrCl3 solutions in 1:1 proportion in the completely discharged state. In the completely charged state, the anolyte comprises CrCl2 and FeCl2 in 1:1 proportion and the catholyte comprises FeCl3 and CrCl3 in 1:1 proportion. In this way, any cross-diffusion of species merely appears as a Coulombic inefficiency, and over time the 1:1 charge balance is maintained. Although the above example describes a Fe/Cr system, it is generally applicable to other Redox couples, such as for example all-Vanadium systems, (see Skyllas-Kazacos in U.S. Pat. No. 4,786,567, incorporated by reference).
One of the major problems of such redox cells is maintaining the charge balance between the anodic and cathodic sides of the cell. If there are no parasitic reactions other than the fundamental redox reactions, then the two sides are always in a charge balanced state. However, in reality parasitic reactions do occur, and after many cycles of charge and discharge, a marked difference with respect to the state of charge of the two electrolyte solutions may develop.
Using the Fe/Cr system as a non limiting example, under ideal conditions (i.e. no parasitic reactions occur) for every Fe3+ ion in the cathode tank there is a Cr2+ ion in the anode tank, and for every Fe2+ ion in the cathode tank, there is a Cr3+ ion in the anode tank. However, in practice, during the charging process, though Fe2+ oxidation proceeds with nearly 100% current yield, reduction of Cr3+ generates hydrogen as a side reaction on the graphite electrodes (see for example U.S. Pat. No. 3,996,064, U.S. Pat. No. 4,382,116 and EPO 0312875), resulting in a higher state of charge of the iron electrolyte i.e. in an excess of Fe3+ ions. Other examples of parasitic reactions include, oxygen (internal or external to the system) oxidizing Fe2+ to Fe3+, or Cr2+ to Cr3+; Cr2+ reducing water to become Cr3+; or during charging, hydrogen generation on the anode in competition with Cr3+ reduction, while Fe2+ oxidation takes place on the cathode.
When this charge imbalance occurs, a rebalancing mechanism is required to return the electrolytes to their charge balanced state. Various rebalancing methods are known to those in the art. Again, using the Fe/Cr system as a non limiting example, charge rebalancing may be achieved by reducing Fe3+ to Fe2+ with hydrogen, (see Thaller in U.S. Pat. No. 4,159,366, herein incorporated by reference).
Before such a rebalancing measure should be taken, it is important to know the state of charge on each side of the cell. Otherwise, an act of unnecessary rebalancing may lead to a worse state of balance and/or waste of energy.
Typically, the state of the charge is determined using separate Open Circuit Voltage (OCV) cell, as described by Hagedorn and Thaller (NASA/TM-81464). The OCV cell is the same as the Redox cell except that there is very high load impedance across the electrodes. Voltage is measured across this resistance, and as the current is virtually zero, it is very close to the OCV, which in turn is directly related to the concentrations of the reactants through Nernst Equation:
  OCV  =            E      .              -        0.0592              ⁢                  ⁢    Ln    ⁢                  ⁢                            [                      Cr                          3              +                                ]                ⁡                  [                      Fe                          2              +                                ]                                      [                      Cr                          2              +                                ]                ⁡                  [                      Fe                          3              +                                ]                    where Eo is the standard potential. Measurement of the OCV is therefore an indirect measurement of the ratio of reactants in the system. However, measuring the OCV in order to determine the charge balance of the redox system has its limitations. For example, it is difficult to differentiate between a system that is out of balance due to parasitic reactions from a partially discharged system. To overcome this limitation one would have to have an accurate coulomb gauge in place, which in turn may be subject to cumulative error after many cycles of operation and the presence of an internal shunt current generated through conductive liquid paths across cells.
Alternatively, it is possible to measure the state of charge of the anode and cathode tanks independently by using cells with respect to standard electrodes, such as Pt/H2 or Ag/AgCl, commonly used in the field of electrochemistry. However, such in situ methods have short life spans, due to the contamination and consumption of the standard electrodes, requiring their frequent replacement. Additionally, cross diffusion of reactants into opposite half-cells renders these measurements unreliable. Further, in premixed solution redox cells, as noted earlier it becomes even more difficult to measure the state of charge and state of balance by just measuring the OCV, as there are many reactants and unknown concentrations involved in the Nernst Equation. Even when reactants are unmixed, cross-diffusion renders the calculation of state of charge and state of balance from OCV erroneous.
It is therefore highly desirable in a redox flow cell system to have a reliable method for determining the state of charge of each electrolyte and the overall state of charge balance, independent of any electrochemical measurement.